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Chemistry 2nd Semester Final Review

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Chemistry 2nd Semester Final Review Powered By Docstoc
					Chemistry 2nd Semester Final Review
Due Thursday, June 17th, 8 am

Gas Laws (Ch. 13):
Subtopics: atmospheric pressure, Dalton’s law of partial pressures, combined gas law, Ideal gas law, gas
stoichiometry (22.4 L/mol for a gas at STP)
Practice questions:
1. If a gas occupies 3.8 L at a pressure of 2.71 atm, what would the pressure be if the volume changes to 1.47 L?
    (Temp. and amt. of gas constant)
2. A gas has a volume of 1.49 L at a temperature of 34.75 °C. What would the volume be at 78.41 °C? (pressure &
    amt. of gas constant)
3. What volume is occupied by 8.47 g of hydrogen gas at 84.7 °C and 1.04 atm?
4. What volume is occupied by 56.75 g of oxygen gas at STP?
5. If 15.71 g of oxygen gas has a volume of 8.14 L, What is the volume of 48.39 g?
6. If hydrogen gas occupies a volume of 2.49 L at 3.41 atm, what would the volume be at 1.37 atm? (Temp. and amt.
    of gas constant)
7. Calcium carbonate decomposes at high temperatures to form carbon dioxide and calcium oxide: CaCO3(s) =>
    CO2(g) + CaO(s). How many grams of calcium carbonate will I need to form 3.45 liters of carbon dioxide?
8. How many liters of water can be made from 55 grams of oxygen gas and an excess of hydrogen at a pressure of
    12.4 atm and a temperature of 85C?
9. How many liters of water can be made from 34 grams of oxygen gas and 6.0 grams of hydrogen gas at STP? What
    is the limiting reactant for this reaction?
10. Define ideal gas.
11. Describe the properties of gases.
12. What is atmospheric pressure? How is it measured? What are the general units used to describe how much force is
    on an object?
13. Write the equations for the following laws and explain what each means: Boyle’s law, Charles’ law, Gay-Lussac’s
    law, Avogadro’s law, combined gas law, ideal gas law

Intermolecular forces, temperature, and heat (Ch. 14):
Subtopics: differences between states of matter, temperature, temperature scales, kinetic energy, specific heat, phase
changes, heating/cooling
14. Describe what happens on a molecular level as a block of ice turns into steam. What is different about the water at
    each phase?
15. Define temperature. How is it related to kinetic energy?
16. Convert 45 °C to Kelvins.
17. Define specific heat. What does it mean if a substance has a high specific heat?
18. How much energy is required to heat 40 g of liquid water from 7 °C to 78 °C?

Solutions (Ch. 15 and some Ch. 2):
Subtopics: mixtures, molarity
19. Distinguish between a mixture and a pure substance.
20. Describe what happens in a solution that has reached equilibrium.
21. Give 2 examples of pure substances.
22. Provide a synonym for “homogenous mixture.”
23. Fill in the following table:
Solute                  Solvent                  Example
solid                   solid                    Brass dissolved in gold to make 14 carat “gold” (alloy)
solid                   liquid
gas                     solid
liquid                  liquid
gas                     liquid
gas                     gas
24. What ingredient in Gatorade acts as an electrolyte?
25. How many moles of PbSO4 are needed to make 200 mL of a 3.7M solution?

26. Calculate the molarity of a solution that contains 1.22 grams of hemoglobin (MW = 68300) in 165 ml of solution.

27. What would adding a solute do to the colligative properties of a solvent?

Acids and Bases (Ch. 16):
Subtopics: physical and chemical properties of acids and bases, Arrhenius and Bronsted-Lowry definitions, conjugate
acid-base pairs, pH, [H3O+], neutralization reactions
28. Describe the properties of acids.
29. Describe the properties of bases.
30. What is the pH range for each of the following: acids, bases, neutral substances?
31. List strong acids and strong bases.
32. Explain each of the following definitions of acids and bases: Arrhenius, Bronsted-Lowry, and Lewis

Label each compound in the equations using the following terms: acid, base, conjugate acid, conjugate base:
33. HNO3         +      H2O <===>             H3O+ +          NO3¯
34. NH3          +      H2O <===>             NH4+ +          OH¯
           -                                +             -2
35. HCO3 + H2O <===>                   H3O +          CO3
36. Write the equation for Kw. Fill in the concentration values for pure water.
37. For each problem (vertical column), fill in the missing values.
Example: for problem #1, you are given the hydrogen ion concentration. You need to find the hydroxide ion
concentration, pH, pOH, and if the solution is acidic, basic, or neutral.
Problem #1               #2     #3 #4             #5             #6      #7          #8
[H+]         7.2 X 10-5                           8.2 X 10-11
[OH-]                                  2 X 10-3                          6.2 X 10-12
pH                       10.5                                                        0.4
pOH                             4.7                              12.4
Acidic/
Basic/
Neutral

Reaction rates (Ch. 17):
Subtopics: requirements for a reaction to occur, activation energy, endothermic/exothermic reactions, factors that
affect reaction rate, catalysts
38. Define activation energy.
39. Draw an energy vs. reaction progress graph for each of the following types of reactions: endothermic, exothermic,
    and catalyzed reaction
40. List the factors that can affect reaction rates. Explain what happens to the reaction rate when these factors are
    changed.
41. What is true of each of the following at equilibrium: rates of forward and reverse reactions, concentrations, and
    amounts of reactants and products?
42. Which variable can change the equilibrium constant?
Equilibrium (Ch. 17):
Subtopics: definition of equilibrium, equilibrium constant (what effects it; writing the equilibrium expression), pressure
and equilibrium, temperature and equilibrium, solubility given Ksp
43. What does it mean if K is high? Low?
44. Write an equilibrium expression for this reaction: A + B  C + D.
45. Write an equilibrium expression for this reaction: 2A3 + B4C2  D2.
46. 3H2(g) + N2(g) ⇄ 2 NH3(g)
       Given that this reaction is exothermic, what direction will the equilibrium shift when the temperature of the
       reaction is decreased?
47. 2 NO2(g) ⇄ N2O4(g)
       If a large quantity of argon is added to the container in which this equilibrium is taking place, in what direction
       will the equilibrium shift?
48. NH4OH(aq) ⇄ NH3(g) + H2O(l)
      In what direction will the equilibrium shift if ammonia is removed from the container as soon as it is produced?
49. 2 BH3(g) ⇄ B2H6(g)
       If this equilibrium is taking place in a piston with a volume of 1 L and I compress it so the final volume is 0.5 L,
       in what direction will the equilibrium shift?
50. H2(g) + Cl2(g) ⇄ 2 HCl(g)
        What direction will the equilibrium shift when the partial pressure of
          hydrogen is increased?
51.     Example: Find the concentration of ions present in calcium fluoride (in water) and the molar solubility.
         CaF2(s) --> Ca+2 + 2 F-       Ksp = 2 x 10 -10
Nuclear chemistry (Ch. 19):
Subtopics: atomic number, mass number, isotopes, types of radiation, balancing a nuclear equation, half-life
52. What is the mass number of argon?
53. What is the atomic number of carbon?
54. Fill out this table regarding the following types of radiation:
Type of radiation Atomic # Mass # Charge Synonym, if any
Alpha particle        2
Beta particle                              -1
Gamma                            0
Neutron               0
Positron                                             Positive electron

55. Neutron bombardment of plutonium -239 yields americium-240 and another
   particle. Write the nuclear equation and identify the other particle produced.
56. When bombarded with neutrons, lithium-6 produces an alpha particle and an
    isotope of hydrogen. Write the nuclear equation for this reaction. What isotope
    of hydrogen is produced?
57. What is the half-life of a 100.0 g sample of nitrogen-16 that decays to 12.5 g of
     nitrogen-16 in 21.6 s?
Organic and Biochemistry (Ch. 20 and 21):
Subtopics: types of carbon-carbon bonding, types of biochemicals
58. Describe the types of bonds that carbon is able to form with itself.
59. Describe the characteristics of each of the main categories of biomolecules.

				
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